Lecture 2
Atoms and Molecules
Atoms

Matter is any substance in the universe that has mass and occupies space

All matter is composed of extremely small particles called atoms

Every atom has the same basic structure
 Core nucleus of protons and neutrons
 Orbiting cloud of electrons
Atoms

Atomic number
 Number of protons

Atomic mass
 Number of protons and
neutrons

Element
 A substance that cannot be
broken down by ordinary
chemical means
Isotopes
 Isotopes are atoms with the same number of protons but
different numbers of neutrons
99% of all
carbon
Different
atomic mass
Same atomic
number
Radioactive Decay
Radioactive isotope dating
 The nucleus of an unstable
isotope breaks down into particles
with lower atomic numbers
 Radioactive isotopes are used in
 1. Medicine
 Tracers are taken up and
used by the body
 Emissions are detected
using special lab
equipment
 2. Dating fossils
 The rate of decay of a
radioactive element is
constant
 The amount of decay can
be used to date fossils
Energy
 The capacity to do work (put matter into motion)
 Types of energy
 Kinetic – energy in action
 Potential – energy of position; stored (inactive) energy
PLAY
Energy Concepts
Electrons
 Electrons have energy due to their relative orbital position
(potential energy)
 Electron shells, or energy levels, surround the nucleus of an atom
 Valence shell – outermost energy level containing chemically active
electrons
 Bonds are formed using the electrons in the outermost energy level
The Octet Rule
 Octet rule – except for the first shell which is full with two
electrons, atoms interact in a manner to have eight electrons
in their valence shell

Inert elements have their outermost
 Reactive elements do not have
energy level fully occupied by electrons
their outermost energy level fully
occupied by electrons
Molecules

A molecule is a group of atoms held together by energy

The holding force is called a chemical bond

There are three kinds of chemical bonds
1. Ionic bonds
2. Covalent bonds
3. Hydrogen bonds
Ionic Bonds




Ionic bonds form between atoms by the transfer of one or more
electrons
Ionic compounds form crystals instead of individual molecules
Example: NaCl (sodium chloride)
Two key properties
1. Strong: But not as strong as covalent bonds
2. Not directional: They are not formed between particular ions in the
compound
Ionic Bonds
Covalent Bonds
 Covalent bonds are formed by the sharing of two or more electrons
 Electron sharing produces molecules
 Two key properties
1. Strong: The strength
increases with the
number of shared
electrons
2. Very directional: They
are formed between
two specific atoms
Water molecules contain two covalent bonds
Comparison of Bonds
 Electrons shared equally
between atoms produce
nonpolar molecules
 Electrons shared unequally
produces polar molecules
 Atoms with six or seven valence
shell electrons are
electronegative
 Atoms with one or two valence
shell electrons are
electropositive
Hydrogen Bonds
 Formed by the attraction of
opposite partial electric charges
between two polar molecules
 Too weak to bind atoms together
 Common in dipoles such as
water
 Responsible for surface tension
in water
 Important as intramolecular
bonds, giving the molecule a
three-dimensional shape
Hydrogen bonding in water molecules
Chemical Reactions
 Occur when chemical bonds are formed, rearranged, or
broken
 Are written in symbolic form using chemical equations
 Chemical equations contain:
 Number and type of reacting substances, and products
produced
 Relative amounts of reactants and products
Patterns of Chemical Reactions
 Combination reactions: Synthesis reactions which always
involve bond formation
A + B  AB
 Decomposition reactions: Molecules are broken down into
smaller molecules
AB  A + B
 Exchange reactions: Bonds are both made and broken
AB + C  AC + B
 All chemical reactions are theoretically reversible
A + B  AB
AB  A + B
 If neither a forward nor reverse reaction is dominant,
chemical equilibrium is reached
Oxidation-Reduction (Redox) Reactions
 Reactants losing electrons are electron donors and are
oxidized
 Reactants taking up electrons are electron acceptors and
become reduced
 Generally, the atom that is reduced contains the most
energy
Energy Flow in Chemical Reactions
 Exergonic reactions – reactions that release energy
 Endergonic reactions – reactions whose products contain
more potential energy than did its reactants
Factors Influencing Rate of Chemical Reactions
 Temperature – chemical reactions proceed quicker at
higher temperatures
 Particle size – the smaller the particle the faster the
chemical reaction
 Concentration – higher reacting particle concentrations
produce faster reactions
 Catalysts – increase the rate of a reaction without being
chemically changed
 Enzymes – biological catalysts
Hydrogen Bonds Give Water Unique Properties
 Water molecules are polar molecules
 They can thus form hydrogen bonds with each other and with other polar
molecules
 Each hydrogen bond is very weak
 However, the cumulative effect of enormous numbers can make them quite
strong
 Hydrogen bonding is responsible for many of the physical properties of
water
 Heat Storage
 A large input of thermal energy is required to disrupt the organization of
liquid water
 This minimizes temperature changes
 Ice Formation
 At low temperatures, hydrogen bonds don’t break
 Water forms a regular crystal structure that floats
 High Heat of Vaporization
 At high temperatures, hydrogen bonds do break
 Water is changed into vapor
PLAY
Water Transport
Hydrogen Bonds Give Water Unique Properties
 Cohesion
 Attraction of water molecules to other
water molecules
 Example: Surface tension
 Adhesion
 Attraction of water molecules to other
polar molecules
 Example: Capillary action
Water strider
Hydrogen Bonds Give Water Unique Properties
 High Polarity
 Polar molecules are termed hydrophilic
 Water-loving
 All polar molecules that dissolve in water are termed
soluble
 Nonpolar molecules are termed hydrophobic
 Water-fearing
 These do not form hydrogen bonds and are therefore
not water soluble
Water Ionizes
 Covalent bonds within a water molecule sometimes break
spontaneously
H 2O
OH–
hydroxide
ion
+
H+
hydrogen
ion
 This process of spontaneous ion formation is called ionization
 It is not common because of the strength of covalent bonds
pH
 A convenient way to
express the hydrogen
ion concentration of a
solution
pH = _ log [H+]
 The pH scale is logarithmic
 A difference of one unit
represents a ten-fold
change in H+ concentration
 Acid: Dissociates in water
to increase H+
concentration
 Base: Combines with H+
when dissolved in water
Buffers
 Hydrogen ion reservoirs that take up or release H+ as needed
 The key buffer in blood is an acid-base pair